Algorithm for composing the electronic formula of an element:
1. Determine the number of electrons in an atom using the Periodic Table of Chemical Elements D.I. Mendeleev.
2. Using the number of the period in which the element is located, determine the number of energy levels; number of electrons on the last electronic level corresponds to the group number.
3. Divide the levels into sublevels and orbitals and fill them with electrons in accordance with the rules for filling orbitals:
It must be remembered that the first level contains a maximum of 2 electrons 1s 2, on the second - a maximum of 8 (two s and six R: 2s 2 2p 6), on the third - a maximum of 18 (two s, six p, and ten d: 3s 2 3p 6 3d 10).
- Principal quantum number n should be minimal.
- First to fill s- sublevel, then р-, d- b f- sublevels.
- Electrons fill the orbitals in order of increasing energy of the orbitals (Klechkovsky's rule).
- Within a sublevel, electrons first occupy free orbitals one by one, and only after that they form pairs (Hund’s rule).
- There cannot be more than two electrons in one orbital (Pauli principle).
Examples.
1. Let's create the electronic formula of nitrogen. IN periodic table nitrogen is at number 7.
2. Let's create the electronic formula for argon. Argon is number 18 on the periodic table.
1s 2 2s 2 2p 6 3s 2 3p 6.
3. Let's create the electronic formula of chromium. Chromium is number 24 on the periodic table.
1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5
Energy diagram of zinc.
4. Let's create the electronic formula of zinc. Zinc is number 30 on the periodic table.
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10
Please note that part of the electronic formula, namely 1s 2 2s 2 2p 6 3s 2 3p 6, is electronic formula argon.
The electronic formula of zinc can be represented as:
Chemicals are what the world around us is made of.
The properties of each chemical substance are divided into two types: chemical, which characterize its ability to form other substances, and physical, which are objectively observed and can be considered in isolation from chemical transformations. For example, the physical properties of a substance are its state of aggregation(solid, liquid or gaseous), thermal conductivity, heat capacity, solubility in different environments(water, alcohol, etc.), density, color, taste, etc.
Transformations of some chemical substances in other substances are called chemical phenomena or chemical reactions. It should be noted that there are also physical phenomena that are obviously accompanied by changes in some physical properties substances without being converted into other substances. TO physical phenomena, for example, include the melting of ice, freezing or evaporation of water, etc.
About what takes place during a process chemical phenomenon, we can conclude by observing characteristic features chemical reactions, such as color change, sedimentation, gas evolution, heat and/or light.
For example, a conclusion about the occurrence of chemical reactions can be made by observing:
Formation of sediment when boiling water, called scale in everyday life;
The release of heat and light when a fire burns;
Change in color of a cut of a fresh apple in air;
Formation of gas bubbles during dough fermentation, etc.
The smallest particles of a substance that undergo virtually no changes during chemical reactions, but only connect with each other in a new way, are called atoms.
The very idea of the existence of such units of matter arose back in ancient Greece in the minds of ancient philosophers, which actually explains the origin of the term “atom”, since “atomos” literally translated from Greek means “indivisible”.
However, contrary to the idea of ancient Greek philosophers, atoms are not the absolute minimum of matter, i.e. they themselves have a complex structure.
Each atom consists of so-called subatomic particles– protons, neutrons and electrons, denoted by the symbols p + , n o and e − , respectively. The superscript in the notation used indicates that the proton has a unit positive charge, electron – single negative charge, but the neutron has no charge.
As for the qualitative structure of an atom, in each atom all protons and neutrons are concentrated in the so-called nucleus, around which the electrons form an electron shell.
The proton and neutron have almost the same masses, i.e. m p ≈ m n, and the mass of the electron is almost 2000 times less than the mass of each of them, i.e. m p /m e ≈ m n /m e ≈ 2000.
Since the fundamental property of an atom is its electrical neutrality, and the charge of one electron equal to charge one proton, from this we can conclude that the number of electrons in any atom is equal to the number of protons.
For example, the table below shows the possible composition of atoms:
Type of atoms with equal charge nuclei, i.e. With the same number protons in their nuclei are called a chemical element. Thus, from the table above we can conclude that atom1 and atom2 belong to one chemical element, and atom3 and atom4 belong to another chemical element.
Each chemical element has its own name and individual symbol, which is read in a certain way. So, for example, the simplest chemical element, the atoms of which contain only one proton in the nucleus, is called “hydrogen” and is denoted by the symbol “H”, which is read as “ash”, and a chemical element with a nuclear charge of +7 (i.e. containing 7 protons) - “nitrogen”, has the symbol “N”, which is read as “en”.
As can be seen from the table above, atoms of one chemical element may differ in the number of neutrons in the nuclei.
Atoms belonging to the same chemical element, but having different quantities neutrons and, as a consequence, mass are called isotopes.
For example, the chemical element hydrogen has three isotopes - 1 H, 2 H and 3 H. The indices 1, 2 and 3 above the symbol H mean the total number of neutrons and protons. Those. Knowing that hydrogen is a chemical element, which is characterized by the fact that there is one proton in the nuclei of its atoms, we can conclude that in the 1 H isotope there are no neutrons at all (1-1 = 0), in the 2 H isotope - 1 neutron (2-1=1) and in the 3 H isotope – two neutrons (3-1=2). Since, as already mentioned, the neutron and proton have the same masses, and the mass of the electron is negligibly small in comparison with them, this means that the 2 H isotope is almost twice as heavy as the 1 H isotope, and the 3 H isotope is even three times heavier . Due to such a large scatter in the masses of hydrogen isotopes, the isotopes 2 H and 3 H were even assigned separate individual names and symbols, which is not typical for any other chemical element. The 2H isotope was named deuterium and given the symbol D, and the 3H isotope was given the name tritium and given the symbol T.
If we take the mass of the proton and neutron as one, and neglect the mass of the electron, in fact the upper left index, in addition to the total number of protons and neutrons in the atom, can be considered its mass, and therefore this index is called the mass number and is designated by the symbol A. Since the charge of the nucleus of any Protons correspond to the atom, and the charge of each proton is conventionally considered equal to +1, the number of protons in the nucleus is called the charge number (Z). By denoting the number of neutrons in an atom as N, the relationship between mass number, charge number and number of neutrons can be expressed mathematically as:
According to modern ideas, the electron has a dual (particle-wave) nature. It has the properties of both a particle and a wave. Like a particle, an electron has mass and charge, but at the same time, the flow of electrons, like a wave, is characterized by the ability to diffraction.
To describe the state of an electron in an atom, the representations are used quantum mechanics, according to which the electron does not have a specific trajectory and can be located at any point in space, but with different probabilities.
The region of space around the nucleus where an electron is most likely to be found is called an atomic orbital.
An atomic orbital can have various shapes, size and orientation. An atomic orbital is also called an electron cloud.
Graphically, one atomic orbital is usually denoted as a square cell:
Quantum mechanics has an extremely complex mathematical apparatus, therefore within school course chemistry, only the consequences of quantum mechanical theory are considered.
According to these consequences, any atomic orbital and the electron located in it are completely characterized by 4 quantum numbers.
- The main quantum number - n - determines total energy electron in a given orbital. Range of values of the main quantum number – all integers, i.e. n = 1,2,3,4, 5, etc.
- The orbital quantum number - l - characterizes the shape of the atomic orbital and can take any integer value from 0 to n-1, where n, recall, is the main quantum number.
Orbitals with l = 0 are called s-orbitals. s-Orbitals are spherical in shape and have no directionality in space:
Orbitals with l = 1 are called p-orbitals. These orbitals have the shape of a three-dimensional figure eight, i.e. a shape obtained by rotating a figure eight around an axis of symmetry, and outwardly resemble a dumbbell:
Orbitals with l = 2 are called d-orbitals, and with l = 3 – f-orbitals. Their structure is much more complex.
3) Magnetic quantum number – m l – determines spatial orientation specific atomic orbital and expresses the projection orbital moment impulse per direction magnetic field. The magnetic quantum number m l corresponds to the orientation of the orbital relative to the direction of the external magnetic field strength vector and can take any integer values from –l to +l, including 0, i.e. total possible values equals (2l+1). So, for example, for l = 0 m l = 0 (one value), for l = 1 m l = -1, 0, +1 (three values), for l = 2 m l = -2, -1, 0, +1 , +2 (five values of magnetic quantum number), etc.
So, for example, p-orbitals, i.e. orbitals with an orbital quantum number l = 1, having the shape of a “three-dimensional figure of eight,” correspond to three values of the magnetic quantum number (-1, 0, +1), which, in turn, correspond to three directions perpendicular to each other in space.
4) The spin quantum number (or simply spin) - m s - can conditionally be considered responsible for the direction of rotation of the electron in the atom; it can take on values. Electrons with different spins are indicated by vertical arrows pointing in different sides: ↓ and .
The set of all orbitals in an atom that have the same principal quantum number is called the energy level or electron shell. Any arbitrary energy level with some number n consists of n 2 orbitals.
Many orbitals with the same values principal quantum number and orbital quantum number represents an energy sublevel.
Each energy level, which corresponds to the principal quantum number n, contains n sublevels. In turn, each energy sublevel with orbital quantum number l consists of (2l+1) orbitals. Thus, the s sublevel consists of one s orbital, the p sublevel consists of three p orbitals, the d sublevel consists of five d orbitals, and the f sublevel consists of seven f orbitals. Since, as already mentioned, one atomic orbital is often denoted by one square cell, the s-, p-, d- and f-sublevels can be graphically represented in the following way:
Each orbital corresponds to an individual strictly specific set three quantum numbers n, l and m l.
The distribution of electrons among orbitals is called the electron configuration.
Filling atomic orbitals electrons occurs in accordance with three conditions:
- Minimum energy principle: Electrons fill orbitals starting from the lowest energy sublevel. The sequence of sublevels in increasing order of their energies is as follows: 1s<2s<2p<3s<3p<4s≤3d<4p<5s≤4d<5p<6s…;
To make it easier to remember this sequence of filling out electronic sublevels, the following graphic illustration is very convenient:
- Pauli principle: Each orbital can contain no more than two electrons.
If there is one electron in an orbital, then it is called unpaired, and if there are two, then they are called an electron pair.
- Hund's rule: the most stable state of an atom is one in which, within one sublevel, the atom has the maximum possible number of unpaired electrons. This most stable state of the atom is called the ground state.
In fact, the above means that, for example, the placement of 1st, 2nd, 3rd and 4th electrons in three orbitals of the p-sublevel will be carried out as follows:
The filling of atomic orbitals from hydrogen, which has a charge number of 1, to krypton (Kr), with a charge number of 36, will be carried out as follows:
Such a representation of the order of filling of atomic orbitals is called an energy diagram. Based on the electronic diagrams of individual elements, it is possible to write down their so-called electronic formulas (configurations). So, for example, an element with 15 protons and, as a consequence, 15 electrons, i.e. phosphorus (P) will have the following energy diagram:
When converted into an electronic formula, the phosphorus atom will take the form:
15 P = 1s 2 2s 2 2p 6 3s 2 3p 3
The normal size numbers to the left of the sublevel symbol show the energy level number, and the superscripts to the right of the sublevel symbol show the number of electrons in the corresponding sublevel.
Below are the electronic formulas of the first 36 elements of the periodic table by D.I. Mendeleev.
| period | Item no. | symbol | Name | electronic formula |
| I | 1 | H | hydrogen | 1s 1 |
| 2 | He | helium | 1s 2 | |
| II | 3 | Li | lithium | 1s 2 2s 1 |
| 4 | Be | beryllium | 1s 2 2s 2 | |
| 5 | B | boron | 1s 2 2s 2 2p 1 | |
| 6 | C | carbon | 1s 2 2s 2 2p 2 | |
| 7 | N | nitrogen | 1s 2 2s 2 2p 3 | |
| 8 | O | oxygen | 1s 2 2s 2 2p 4 | |
| 9 | F | fluorine | 1s 2 2s 2 2p 5 | |
| 10 | Ne | neon | 1s 2 2s 2 2p 6 | |
| III | 11 | Na | sodium | 1s 2 2s 2 2p 6 3s 1 |
| 12 | Mg | magnesium | 1s 2 2s 2 2p 6 3s 2 | |
| 13 | Al | aluminum | 1s 2 2s 2 2p 6 3s 2 3p 1 | |
| 14 | Si | silicon | 1s 2 2s 2 2p 6 3s 2 3p 2 | |
| 15 | P | phosphorus | 1s 2 2s 2 2p 6 3s 2 3p 3 | |
| 16 | S | sulfur | 1s 2 2s 2 2p 6 3s 2 3p 4 | |
| 17 | Cl | chlorine | 1s 2 2s 2 2p 6 3s 2 3p 5 | |
| 18 | Ar | argon | 1s 2 2s 2 2p 6 3s 2 3p 6 | |
| IV | 19 | K | potassium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 |
| 20 | Ca | calcium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 | |
| 21 | Sc | scandium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 | |
| 22 | Ti | titanium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 | |
| 23 | V | vanadium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 | |
| 24 | Cr | chromium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 here we observe the jump of one electron with s on d sublevel | |
| 25 | Mn | manganese | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 | |
| 26 | Fe | iron | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 | |
| 27 | Co | cobalt | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 | |
| 28 | Ni | nickel | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 | |
| 29 | Cu | copper | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 here we observe the jump of one electron with s on d sublevel | |
| 30 | Zn | zinc | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 | |
| 31 | Ga | gallium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 1 | |
| 32 | Ge | germanium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 2 | |
| 33 | As | arsenic | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 | |
| 34 | Se | selenium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 4 | |
| 35 | Br | bromine | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 | |
| 36 | Kr | krypton | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 |
As already mentioned, in their ground state, electrons in atomic orbitals are located according to the principle of least energy. However, in the presence of empty p-orbitals in the ground state of the atom, often, by imparting excess energy to it, the atom can be transferred to the so-called excited state. For example, a boron atom in its ground state has an electronic configuration and an energy diagram of the following form:
5 B = 1s 2 2s 2 2p 1
And in an excited state (*), i.e. When some energy is imparted to a boron atom, its electron configuration and energy diagram will look like this:
5 B* = 1s 2 2s 1 2p 2
Depending on which sublevel in the atom is filled last, chemical elements are divided into s, p, d or f.
Finding s, p, d and f elements in the table D.I. Mendeleev:
- The s-elements have the last s-sublevel to be filled. These elements include elements of the main (on the left in the table cell) subgroups of groups I and II.
- For p-elements, the p-sublevel is filled. The p-elements include the last six elements of each period, except the first and seventh, as well as elements of the main subgroups of groups III-VIII.
- d-elements are located between s- and p-elements in large periods.
- f-Elements are called lanthanides and actinides. They are listed at the bottom of the D.I. table. Mendeleev.
Electrons
The concept of atom arose in the ancient world to designate particles of matter. Translated from Greek, atom means “indivisible.”
The Irish physicist Stoney, based on experiments, came to the conclusion that electricity is carried by the smallest particles existing in the atoms of all chemical elements. In 1891, Stoney proposed to call these particles electrons, which means “amber” in Greek. A few years after the electron got its name, the English physicist Joseph Thomson and the French physicist Jean Perrin proved that electrons carry a negative charge. This is the smallest negative charge, which in chemistry is taken as one (-1). Thomson even managed to determine the speed of the electron (the speed of the electron in the orbit is inversely proportional to the orbit number n. The radii of the orbits increase in proportion to the square of the orbit number. In the first orbit of the hydrogen atom (n=1; Z=1) the speed is ≈ 2.2·106 m/ s, that is, about a hundred times less than the speed of light c = 3·108 m/s) and the mass of the electron (it is almost 2000 times less than the mass of the hydrogen atom).
State of electrons in an atom
The state of an electron in an atom is understood as a set of information about the energy of a particular electron and the space in which it is located. An electron in an atom does not have a trajectory of motion, i.e. we can only talk about the probability of finding it in the space around the nucleus.
It can be located in any part of this space surrounding the nucleus, and the totality of its various positions is considered as an electron cloud with a certain negative charge density. Figuratively, this can be imagined this way: if it were possible to photograph the position of an electron in an atom after hundredths or millionths of a second, as in a photo finish, then the electron in such photographs would be represented as dots. If countless such photographs were superimposed, the picture would be of an electron cloud with the greatest density where there would be the most of these points.
The space around the atomic nucleus in which an electron is most likely to be found is called an orbital. It contains approximately 90% electronic cloud, and this means that about 90% of the time the electron is in this part of space. They are distinguished by shape 4 currently known types of orbitals, which are designated by Latin letters s, p, d and f. A graphical representation of some forms of electron orbitals is presented in the figure.
The most important characteristic of the motion of an electron in a certain orbital is energy of its connection with the nucleus. Electrons with similar energy values form a single electron layer, or energy level. Energy levels are numbered starting from the nucleus - 1, 2, 3, 4, 5, 6 and 7.
The integer n, indicating the number of the energy level, is called the principal quantum number. It characterizes the energy of electrons occupying a given energy level. Electrons of the first energy level, closest to the nucleus, have the lowest energy. Compared to electrons of the first level, electrons of subsequent levels will be characterized by a large supply of energy. Consequently, the electrons of the outer level are least tightly bound to the atomic nucleus.
The largest number of electrons at an energy level is determined by the formula:
N = 2n 2 ,
where N is the maximum number of electrons; n is the level number, or the main quantum number. Consequently, at the first energy level closest to the nucleus there can be no more than two electrons; on the second - no more than 8; on the third - no more than 18; on the fourth - no more than 32.
Starting from the second energy level (n = 2), each of the levels is divided into sublevels (sublayers), slightly different from each other in the binding energy with the nucleus. The number of sublevels is equal to the value of the main quantum number: the first energy level has one sublevel; the second - two; third - three; fourth - four sublevels. The sublevels, in turn, are formed by orbitals. Each valuen corresponds to the number of orbitals equal to n.
Sublevels are usually denoted by Latin letters, as well as the shape of the orbitals of which they consist: s, p, d, f.
Protons and Neutrons
An atom of any chemical element is comparable to a tiny solar system. Therefore, this model of the atom, proposed by E. Rutherford, is called planetary.
The atomic nucleus, in which the entire mass of the atom is concentrated, consists of particles of two types - protons and neutrons.
Protons have a charge equal to the charge of electrons, but opposite in sign (+1), and a mass equal to the mass of a hydrogen atom (it is taken as one in chemistry). Neutrons carry no charge, they are neutral and have a mass equal to the mass of a proton.
Protons and neutrons together are called nucleons (from the Latin nucleus - nucleus). The sum of the number of protons and neutrons in an atom is called the mass number. For example, the mass number of an aluminum atom is:
13 + 14 = 27
number of protons 13, number of neutrons 14, mass number 27
Since the mass of the electron, which is negligibly small, can be neglected, it is obvious that the entire mass of the atom is concentrated in the nucleus. Electrons are designated e - .
Since the atom electrically neutral, then it is also obvious that the number of protons and electrons in an atom is the same. It is equal to the serial number of the chemical element assigned to it in the Periodic Table. The mass of an atom consists of the mass of protons and neutrons. Knowing the atomic number of the element (Z), i.e. the number of protons, and the mass number (A), equal to the sum of the numbers of protons and neutrons, you can find the number of neutrons (N) using the formula:
N = A - Z
For example, the number of neutrons in an iron atom is:
56 — 26 = 30
Isotopes
Varieties of atoms of the same element that have the same nuclear charge but different mass numbers are called isotopes. Chemical elements found in nature are a mixture of isotopes. Thus, carbon has three isotopes with masses 12, 13, 14; oxygen - three isotopes with masses 16, 17, 18, etc. The relative atomic mass of a chemical element usually given in the Periodic Table is the average value of the atomic masses of a natural mixture of isotopes of a given element, taking into account their relative abundance in nature. The chemical properties of isotopes of most chemical elements are exactly the same. However, hydrogen isotopes vary greatly in properties due to the dramatic multiple increase in their relative atomic mass; they are even given individual names and chemical symbols.
Elements of the first period
Diagram of the electronic structure of the hydrogen atom:
Diagrams of the electronic structure of atoms show the distribution of electrons across electronic layers (energy levels).
Graphic electronic formula of the hydrogen atom (shows the distribution of electrons by energy levels and sublevels):
Graphic electronic formulas of atoms show the distribution of electrons not only among levels and sublevels, but also among orbitals.
In a helium atom, the first electron layer is complete - it has 2 electrons. Hydrogen and helium are s-elements; The s-orbital of these atoms is filled with electrons.
For all elements of the second period the first electronic layer is filled, and electrons fill the s- and p-orbitals of the second electron layer in accordance with the principle of least energy (first s and then p) and the Pauli and Hund rules.
In the neon atom, the second electron layer is complete - it has 8 electrons.
For atoms of elements of the third period, the first and second electronic layers are completed, so the third electronic layer is filled, in which electrons can occupy the 3s-, 3p- and 3d-sublevels.
The magnesium atom completes its 3s electron orbital. Na and Mg are s-elements.
In aluminum and subsequent elements, the 3p sublevel is filled with electrons.
Elements of the third period have unfilled 3d orbitals.
All elements from Al to Ar are p-elements. The s- and p-elements form the main subgroups in the Periodic Table.
Elements of the fourth - seventh periods
A fourth electron layer appears in potassium and calcium atoms, and the 4s sublevel is filled, since it has lower energy than the 3d sublevel.
K, Ca - s-elements included in the main subgroups. For atoms from Sc to Zn, the 3d sublevel is filled with electrons. These are 3d elements. They are included in secondary subgroups, their outermost electronic layer is filled, and they are classified as transition elements.
Pay attention to the structure of the electronic shells of chromium and copper atoms. In them, one electron “fails” from the 4s to the 3d sublevel, which is explained by the greater energy stability of the resulting electronic configurations 3d 5 and 3d 10:
In the zinc atom, the third electron layer is complete - all sublevels 3s, 3p and 3d are filled in it, with a total of 18 electrons. In the elements following zinc, the fourth electron layer, the 4p sublevel, continues to be filled.
Elements from Ga to Kr are p-elements.
The krypton atom has an outer layer (fourth) that is complete and has 8 electrons. But there can be a total of 32 electrons in the fourth electron layer; the krypton atom still has unfilled 4d and 4f sublevels. For elements of the fifth period, sublevels are being filled in the following order: 5s - 4d - 5p. And there are also exceptions related to “ failure» electrons, y 41 Nb, 42 Mo, 44 Ru, 45 Rh, 46 Pd, 47 Ag.
In the sixth and seventh periods, f-elements appear, i.e., elements in which the 4f- and 5f-sublevels of the third outside electronic layer are filled, respectively.
4f elements are called lanthanides.
5f elements are called actinides.
The order of filling electronic sublevels in the atoms of elements of the sixth period: 55 Cs and 56 Ba - 6s elements; 57 La … 6s 2 5d x - 5d element; 58 Ce - 71 Lu - 4f elements; 72 Hf - 80 Hg - 5d elements; 81 T1 - 86 Rn - 6d elements. But here, too, there are elements in which the order of filling the electronic orbitals is “violated,” which, for example, is associated with the greater energy stability of half and fully filled f-sublevels, i.e. nf 7 and nf 14. Depending on which sublevel of the atom is filled with electrons last, all elements are divided into four electron families, or blocks:
- s-elements. The s-sublevel of the outer level of the atom is filled with electrons; s-elements include hydrogen, helium and elements of the main subgroups of groups I and II.
- p-elements. The p-sublevel of the outer level of the atom is filled with electrons; p-elements include elements of the main subgroups of groups III-VIII.
- d-elements. The d-sublevel of the pre-external level of the atom is filled with electrons; d-elements include elements of secondary subgroups of groups I-VIII, i.e. elements of plug-in decades of large periods located between s- and p-elements. They are also called transition elements.
- f-elements. The f-sublevel of the third outer level of the atom is filled with electrons; these include lanthanides and antinoids.
The Swiss physicist W. Pauli in 1925 established that in an atom in one orbital there can be no more than two electrons having opposite (antiparallel) spins (translated from English as “spindle”), i.e., having such properties that conditionally can be imagined as the rotation of an electron around its imaginary axis: clockwise or counterclockwise.
This principle is called Pauli principle. If there is one electron in the orbital, then it is called unpaired; if there are two, then these are paired electrons, i.e. electrons with opposite spins. The figure shows a diagram of the division of energy levels into sublevels and the order in which they are filled.
Very often, the structure of the electronic shells of atoms is depicted using energy or quantum cells - so-called graphical electronic formulas are written. For this notation, the following notation is used: each quantum cell is designated by a cell that corresponds to one orbital; Each electron is indicated by an arrow corresponding to the spin direction. When writing a graphical electronic formula, you should remember two rules: Pauli's principle and F. Hund's rule, according to which electrons occupy free cells first one at a time and have the same spin value, and only then pair, but the spins, according to the Pauli principle, will already be oppositely directed.
Hund's rule and Pauli's principle
Hund's rule- a rule of quantum chemistry that determines the order of filling the orbitals of a certain sublayer and is formulated as follows: the total value of the spin quantum number of electrons of a given sublayer must be maximum. Formulated by Friedrich Hund in 1925.
This means that in each of the orbitals of the sublayer, one electron is filled first, and only after the unfilled orbitals are exhausted, a second electron is added to this orbital. In this case, in one orbital there are two electrons with half-integer spins of the opposite sign, which pair (form a two-electron cloud) and, as a result, the total spin of the orbital becomes equal to zero.
Another wording: Lower in energy lies the atomic term for which two conditions are satisfied.
- Multiplicity is maximum
- When the multiplicities coincide, the total orbital momentum L is maximum.
Let us analyze this rule using the example of filling p-sublevel orbitals p-elements of the second period (that is, from boron to neon (in the diagram below, horizontal lines indicate orbitals, vertical arrows indicate electrons, and the direction of the arrow indicates the spin orientation).
Klechkovsky's rule
Klechkovsky's rule - as the total number of electrons in atoms increases (with an increase in the charges of their nuclei, or the serial numbers of chemical elements), atomic orbitals are populated in such a way that the appearance of electrons in an orbital with a higher energy depends only on the main quantum number n and does not depend on all other quantum numbers numbers, including from l. Physically, this means that in a hydrogen-like atom (in the absence of interelectron repulsion), the orbital energy of an electron is determined only by the spatial distance of the electron charge density from the nucleus and does not depend on the characteristics of its motion in the field of the nucleus.
The empirical Klechkovsky rule and the ordering scheme that follows from it are somewhat contradictory to the real energy sequence of atomic orbitals only in two similar cases: for atoms Cr, Cu, Nb, Mo, Ru, Rh, Pd, Ag, Pt, Au, there is a “failure” of an electron with s -sublevel of the outer layer is replaced by the d-sublevel of the previous layer, which leads to an energetically more stable state of the atom, namely: after filling orbital 6 with two electrons s
Let's look at how an atom is built. Keep in mind that we will talk exclusively about models. In practice, atoms are a much more complex structure. But thanks to modern developments, we are able to explain and even successfully predict properties (even if not all). So what is the structure of an atom? What is it “made” of?
Planetary model of the atom
It was first proposed by the Danish physicist N. Bohr in 1913. This is the first theory of atomic structure based on scientific facts. In addition, it laid the foundation for modern thematic terminology. In it, electrons-particles produce rotational movements around the atom according to the same principle as the planets around the Sun. Bohr suggested that they could exist exclusively in orbits located at a strictly defined distance from the nucleus. The scientist could not explain why this was so, from a scientific standpoint, but such a model was confirmed by many experiments. Integer numbers were used to designate orbits, starting with one, which was numbered closest to the nucleus. All these orbits are also called levels. The hydrogen atom has only one level, on which one electron rotates. But complex atoms also have levels. They are divided into components that combine electrons with similar energy potential. So, the second already has two sublevels - 2s and 2p. The third already has three - 3s, 3p and 3d. And so on. First, the sublevels closer to the core are “populated,” and then the distant ones. Each of them can only hold a certain number of electrons. But this is not the end. Each sublevel is divided into orbitals. Let's make a comparison with ordinary life. The electron cloud of an atom is comparable to a city. Levels are streets. Sublevel - private house or apartment. Orbital - room. Each of them “lives” one or two electrons. They all have specific addresses. This was the first diagram of the structure of the atom. And finally, about the addresses of electrons: they are determined by sets of numbers that are called “quantum”.
Wave model of the atom
But over time, the planetary model was revised. A second theory of atomic structure was proposed. It is more advanced and allows you to explain the results of practical experiments. The first one was replaced by the wave model of the atom, which was proposed by E. Schrödinger. Then it was already established that an electron can manifest itself not only as a particle, but also as a wave. What did Schrödinger do? He applied an equation that describes the motion of a wave in Thus, one can find not the trajectory of an electron in an atom, but the probability of its detection at a certain point. What unites both theories is that elementary particles are located at specific levels, sublevels and orbitals. This is where the similarity between the models ends. Let me give you one example: in wave theory, an orbital is a region where an electron can be found with a 95% probability. The rest of the space accounts for 5%. But in the end it turned out that the structural features of atoms are depicted using the wave model, despite the fact that the terminology used is common.
The concept of probability in this case
Why was this term used? Heisenberg formulated the uncertainty principle in 1927, which is now used to describe the movement of microparticles. It is based on their fundamental difference from ordinary physical bodies. What is it? Classical mechanics assumed that a person could observe phenomena without influencing them (observation of celestial bodies). Based on the data obtained, it is possible to calculate where the object will be at a certain point in time. But in the microcosm things are necessarily different. So, for example, it is now not possible to observe an electron without influencing it due to the fact that the energies of the instrument and the particle are incomparable. This leads to changes in its location of the elementary particle, state, direction, speed of movement and other parameters. And it makes no sense to talk about exact characteristics. The uncertainty principle itself tells us that it is impossible to calculate the exact trajectory of an electron around the nucleus. You can only indicate the probability of finding a particle in a certain area of space. This is the peculiarity of the structure of atoms of chemical elements. But this should be taken into account exclusively by scientists in practical experiments.
Atomic composition
But let's concentrate on the entire subject matter. So, in addition to the well-considered electron shell, the second component of the atom is the nucleus. It consists of positively charged protons and neutral neutrons. We are all familiar with the periodic table. The number of each element corresponds to the number of protons it contains. The number of neutrons is equal to the difference between the mass of an atom and its number of protons. There may be deviations from this rule. Then they say that an isotope of the element is present. The structure of an atom is such that it is “surrounded” by an electron shell. usually equals the number of protons. The mass of the latter is approximately 1840 times greater than that of the former, and is approximately equal to the weight of the neutron. The radius of the nucleus is about 1/200,000 the diameter of the atom. It itself has a spherical shape. This, in general, is the structure of the atoms of chemical elements. Despite the difference in mass and properties, they look approximately the same.
Orbits
When talking about what an atomic structure diagram is, one cannot remain silent about them. So, there are these types:
- s. They have a spherical shape.
- p. They look like three-dimensional figure eights or a spindle.
- d and f. They have a complex shape that is difficult to describe in formal language.
An electron of each type can be found with a 95% probability in the corresponding orbital. The information presented must be treated calmly, since it is rather an abstract mathematical model than a physical reality. But with all this, it has good predictive power regarding the chemical properties of atoms and even molecules. The further a level is located from the nucleus, the more electrons can be placed on it. Thus, the number of orbitals can be calculated using a special formula: x 2. Here x is equal to the number of levels. And since up to two electrons can be placed in an orbital, ultimately the formula for their numerical search will look like this: 2x 2.
Orbits: technical data
If we talk about the structure of the fluorine atom, it will have three orbitals. They will all be filled. The energy of orbitals within one sublevel is the same. To designate them, add the layer number: 2s, 4p, 6d. Let's return to the conversation about the structure of the fluorine atom. It will have two s- and one p-sublevel. It has nine protons and the same number of electrons. First one s-level. That's two electrons. Then the second s-level. Two more electrons. And 5 fills the p-level. This is his structure. After reading the following subheading, you can do the necessary steps yourself and make sure of this. If we talk about which fluorine also belongs, it should be noted that they, although in the same group, are completely different in their characteristics. Thus, their boiling point ranges from -188 to 309 degrees Celsius. So why were they united? All thanks to chemical properties. All halogens, and fluorine to the greatest extent, have the highest oxidizing ability. They react with metals and can spontaneously ignite at room temperature without any problems.
How are orbits filled?
By what rules and principles are electrons arranged? We suggest that you familiarize yourself with the three main ones, the wording of which has been simplified for better understanding:
- Principle of least energy. Electrons tend to fill orbitals in order of increasing energy.
- Pauli's principle. One orbital cannot contain more than two electrons.
- Hund's rule. Within one sublevel, electrons first fill empty orbitals, and only then form pairs.
The structure of the atom will help in filling it out and in this case it will become more understandable in terms of image. Therefore, when working practically with the construction of circuit diagrams, it is necessary to keep it at hand.
Example
In order to summarize everything that has been said within the framework of the article, you can draw up a sample of how the electrons of an atom are distributed among their levels, sublevels and orbitals (that is, what the configuration of levels is). It can be depicted as a formula, an energy diagram, or a layer diagram. There are very good illustrations here, which, upon careful examination, help to understand the structure of the atom. So, the first level is filled in first. It has only one sublevel, in which there is only one orbital. All levels are filled sequentially, starting with the smallest. First, within one sublevel, one electron is placed in each orbital. Then pairs are created. And if there are free ones, a switch to another filling subject occurs. And now you can find out for yourself what the structure of the nitrogen or fluorine atom is (which was considered earlier). It may be a little difficult at first, but you can use the pictures to guide you. For clarity, let's look at the structure of the nitrogen atom. It has 7 protons (together with neutrons that make up the nucleus) and the same number of electrons (which make up the electron shell). The first s-level is filled in first. It has 2 electrons. Then comes the second s-level. It also has 2 electrons. And the other three are placed on the p-level, where each of them occupies one orbital.
Conclusion
As you can see, the structure of the atom is not such a difficult topic (if you approach it from the perspective of a school chemistry course, of course). And understanding this topic is not difficult. Finally, I would like to tell you about some features. For example, speaking about the structure of the oxygen atom, we know that it has eight protons and 8-10 neutrons. And since everything in nature tends to balance, two oxygen atoms form a molecule, where two unpaired electrons form a covalent bond. Another stable oxygen molecule, ozone (O3), is formed in a similar way. Knowing the structure of the oxygen atom, you can correctly draw up formulas for oxidative reactions in which the most common substance on Earth participates.
Any substance is made up of very small particles called atoms . An atom is the smallest particle of a chemical element that retains all its characteristic properties. To imagine the size of an atom, it is enough to say that if they could be placed close to one another, then one million atoms would occupy a distance of only 0.1 mm.
Further development of the science of the structure of matter showed that the atom also has a complex structure and consists of electrons and protons. This is how the electronic theory of the structure of matter arose.
In ancient times it was discovered that there are two types of electricity: positive and negative. The amount of electricity contained in the body came to be called charge. Depending on the type of electricity a body possesses, the charge can be positive or negative.
It was also established experimentally that like charges repel, and unlike charges attract.
Let's consider electronic structure of the atom. Atoms are made up of even smaller particles than themselves, called electrons.
DEFINITION:An electron is the smallest particle of matter that has the smallest negative electrical charge.
Electrons orbit around a central nucleus consisting of one or more protons And neutrons, in concentric orbits. Electrons are negatively charged particles, protons are positively charged, and neutrons are neutral (Figure 1.1).
DEFINITION:A proton is the smallest particle of matter that has the smallest positive electrical charge.
The existence of electrons and protons is beyond doubt. Scientists not only determined the mass, charge and size of electrons and protons, but even made them work in various electrical and radio engineering devices.
It was also found that the mass of an electron depends on the speed of its movement and that the electron not only moves forward in space, but also rotates around its axis.
The simplest in structure is the hydrogen atom (Fig. 1.1). It consists of a proton nucleus and an electron rotating at great speed around the nucleus, forming the outer shell (orbit) of the atom. More complex atoms have several shells through which electrons rotate.
These shells are filled with electrons sequentially from the nucleus (Figure 1.2).
Now let's look at it . The outermost shell is called valence, and the number of electrons contained in it is called valence. The farther from the core valence shell, therefore, the less force of attraction each valence electron experiences from the nucleus. Thus, the atom increases the ability to attach electrons to itself in the event that the valence shell is not filled and is located far from the nucleus, or to lose them.
The outer shell electrons can receive energy. If the electrons located in the valence shell receive the required level of energy from external forces, they can break away from it and leave the atom, that is, become free electrons. Free electrons are able to move randomly from one atom to atom. Those materials that contain a large number of free electrons are called conductors
.
Insulators
, is the opposite of conductors. They prevent the flow of electric current. Insulators are stable because the valence electrons of some atoms fill the valence shells of other atoms, joining them. This prevents the formation of free electrons.
Occupy an intermediate position between insulators and conductors semiconductors
, but we'll talk about them later
Let's consider properties of the atom. An atom that has the same number of electrons and protons is electrically neutral. An atom that gains one or more electrons becomes negatively charged and is called a negative ion. If an atom loses one or more electrons, it becomes a positive ion, that is, it becomes positively charged.